化学专业英语之非金属元素 THE NONMETAL ELEMENTS
We noted earlier. that -nonmetals exhibit properties that are greatly different from those of the metals. As a rule, the nonmetals are poor
conductors of electricity (graphitic carbon is an exception) and heat; they are brittle, are often intensely colored, and show an unusually wide range of melting and boiling points. Their molecular structures, usually involving ordinary covalent bonds, vary from the simple diatomic
molecules of H2, Cl2, I2, and N2 to the giant molecules of diamond, silicon and boron.
The nonmetals that are gases at room temperature are the low-molecular weight diatomic molecules and the noble gases that exert very small intermolecular forces. As the molecular weight increases, we encounter a liquid (Br2) and a solid (I2) whose vapor pressures also indicate small intermolecular forces. Certain properties of a few nonmetals are listed in Table 2.
Table 2- Molecular Weights and Melting Points of Certain Nonmetals Molecular Diatomic Molecules Weight H2 N2 F2 O2 Cl2 Br2 I2
Simple diatomic molecules are not formed by the heavier members of
2 28 38 32 71 160 254 Melting Point °C -239.1' -210 -223 -218 -102 -7.3 113 Color None None Pale yellow Pale blue Yellow — green Red — brown Gray—black Groups V and VI at ordinary conditions. This is in direct contrast to the first members of these groups, N2 and O2. The difference arises because of the lower stability of π bonds formed from p orbitals of the third and higher main energy levels as opposed to the second main energy level2. The larger atomic radii and more dense electron clouds of elements of the third period and higher do not allow good parallel overlap of p orbitals necessary for a strong π bond. This is a general phenomenon — strong π bonds are formed only between elements of the second period. Thus, elemental nitrogen and oxygen form stable molecules with both σ and π bonds, but other members of their groups form more stable structures based on σ bonds only at ordinary conditions. Note3 that Group VII elements form diatomic molecules, but π bonds are not required for saturation of valence.
Sulfur exhibits allotropic forms. Solid sulfur exists in two crystalline forms and in an amorphous form. Rhombic sulfur is obtained by
crystallization from a suitable solution, such as CS2, and it melts at 112°C. Monoclinic sulfur is formed by cooling melted sulfur and it melts at 119°C. Both forms of crystalline sulfur melt into S-gamma, which is composed of S8 molecules. The S8 molecules are puckered rings and survive heating to about 160°C. Above 160°C, the S8 rings break open, and some of these fragments combine with each other to form a highly viscous mixture of irregularly shaped coils. At a range of higher
temperatures the liquid sulfur becomes so viscous that it will not pour from its container. The color also changes from straw yellow at sulfur's melting point to a deep reddish-brown as it becomes more viscous. As4 the boiling point of 444 °C is approached, the large-coiled
molecules of sulfur are partially degraded and the liquid sulfur decreases in viscosity. If the hot liquid sulfur is quenched by pouring it into cold water, the amorphous form of sulfur is produced. The structure of
amorphous sulfur consists of large-coiled helices with eight sulfur atoms to each turn of the helix; the overall nature of amorphous sulfur is
described as3 rubbery because it stretches much like ordinary rubber. In a few hours the amorphous sulfur reverts to small rhombic crystals and its rubbery property disappears.
Sulfur, an important raw material in industrial chemistry, occurs as the free element, as SO2 in volcanic regions, as H2S in mineral waters, and in a variety of sulfide ores such as iron pyrite FeS2, zinc blende ZnS, galena PbS and such, and in common formations of gypsum CaSO4 ? 2H2O, anhydrite CaSO4, and barytes BaSO4 ? 2H2O. Sulfur, in one form or another, is used in large quantities for making sulfuric acid, fertilizers, insecticides, and paper.
Sulfur in the form of SO2 obtained in the roasting of sulfide ores is
recovered and converted to sulfuric acid, although in previous years much of this SO2 was discarded through exceptionally tall smokestacks.
Fortunately, it is now economically favorable to recover these gases, thus greatly reducing this type of atmospheric pollution. A typical roasting reaction involves the change:
2 ZnS + 3 O2—2 ZnO + 2 SO2
Phosphorus, below 800℃ consists of tetratomic molecules, P4. Its
molecular structure provides for a covalence of three, as may be expected from the three unpaired p electrons in its atomic structure, and each atom is attached to three others6. Instead of a strictly orthogonal orientation, with the three bonds 90° to each other, the bond angles are only 60°. This supposedly strained structure is stabilized by the mutual interaction of the four atoms (each atom is bonded to the other three), but it is chemically the most active form of phosphorus. This form of phosphorus, the white modification, is spontaneously combustible in air. When heated to 260°C it changes to red phosphorus, whose structure is obscure. Red phosphorus is stable in air but, like all forms of phosphorus, it should be handled carefully because of its tendency to migrate to the bones when ingested, resulting in serious physiological damage.
Elemental carbon exists in one of two crystalline structures — diamond and graphite. The diamond structure, based on tetrahedral bonding of hybridized sp3 orbitals, is encountered among Group IV elements. We may expect that as the bond length increases, the hardness of the diamond-type crystal decreases. Although the tetrahedral structure
persists among the elements in this group — carbon, silicon, germanium,
化学专业英语之非金属元素
